a chemical element, first in the ordering sequence in Mendeleev’s periodic system. Atomic mass, 1.00797. Under ordinary conditions, a gaseous, colorless, odorless, and tasteless element.
History. The treatises of 16th- and 17th-century chemists repeatedly mention the evolution of a flammable gas in the action of acids on metals. In 1766, H. Cavendish collected and studied this evolved gas, naming it “inflammable air.” An advocate of the phlogiston theory, Cavendish assumed that this gas was pure phlogiston. In 1783, A. L. Lavoisier, by analysis and synthesis of water, showed its complex composition. In 1787 he established that “inflammable air” was a new chemical element (H) and gave it its modern name hydrogen, which is Greek for “maker of water.” The root of the word is used in the names of hydrogen compounds and the processes in which hydrogen participates (for example, hydrides, hydrogenation). The present-day Russian name vodorod was proposed by M. F. Solov’ev in 1824.
Hydrogen in nature. Hydrogen is widespread in nature. Its content in the earth’s crust (lithosphere and hydrosphere) is 1 percent by mass and 16 percent on an atomic basis. Hydrogen is a constituent of the substance most abundant on earth—water (11.19 percent H by mass). As a constituent of compounds, hydrogen is present in coal, petroleum, natural gases, clay, and animal and plant organisms (that is, in the composition of proteins, nucleic acids, fats, carbohydrates, and so forth). Hydrogen is rarely found in a free state; small quantities are found in volcanic and other natural gases. Insignificant amounts of free hydrogen (0.0001 percent on an atomic basis) are contained in the atmosphere. As a stream of protons, in the space around the earth, hydrogen forms the inner (proton) Van Allen radiation belt. Hydrogen is the most abundant element in space. In the form of plasma it constitutes about half of the sun’s mass and the masses of most stars and the main part of interstellar medium gases and gaseous nebulas. Hydrogen is found in the atmosphere of some planets and in comets in the form of free H2, methane CH4, ammonia NH3, water H2O, and such radicals as CH, NH, OH, SiH, and PH. In the form of a stream of protons, hydrogen is present in the corpuscular radiation of the sun and in cosmic rays.
Isotopes, atom, and molecule. Ordinary hydrogen is a mixture of two stable isotopes: light hydrogen, or protium (1H), and heavy hydrogen, or deuterium (2H, or D). In natural compounds, there are on the average 6,800 atoms of 1H for every 2H atom. An artificially produced radioactive isotope is superheavy hydrogen, or tritium (3H, or T), with a soft β-radiation and a half-life of 12.262 years. In nature tritium is formed, for example, from atomic nitrogen by the action of cosmic-ray neutrons. The amount of tritium in the atmosphere is insignificant (4 × 10−15 percent of all the hydrogen atoms). A very unstable isotope, 4H, has also been obtained. The mass numbers of the isotopes 1H, 2H, 3H, and 4H, which are, respectively, 1, 2, 3, and 4, indicate that the nucleus of the protium atom contains only one proton; the deuterium, one proton and one neutron; the tritium, one proton and two neutrons; and the 4H, one proton and three neutrons. The great differences in the masses of hydrogen isotopes makes the differences in their physical and chemical properties more noticeable than in the isotopes of other elements. Of the atoms of all elements, the hydrogen atom has the simplest structure; it consists of one nucleus and one electron. The binding energy between the electron and nucleus (ionization potential) is 13.595 electron volts (eV). A neutral hydrogen atom can accept a second electron and form a negative ion H−, where the binding energy of the second electron and a neutral atom (the affinity to an electron) is 0.78 eV. Quantum mechanics allows us to calculate all the possible energy levels of the hydrogen atom and hence to give a full interpretation of its atomic spectrum. The hydrogen atom is used as a model in the quantum mechanical calculations of the energy levels of other, more complex, atoms. The hydrogen molecule, H2, has two atoms united by a covalent chemical bond. The dissociation energy (that is, separation into atoms) is 4.77 eV (1 eV = 1.60210 × 10−19 joule). The equilibrium interatomic distance is 0.7414 Å. At high temperatures, molecular hydrogen dissociates into atoms (the degree of dissociation at 2000° C is 0.0013 and at 5000° C, 0.95). Atomic hydrogen also forms in various chemical reactions (for example, when zinc reacts with hydrochloric acid). However, the existence of hydrogen in the atomic state is very short, and the atoms are recombined into H2 molecules.
Physical and chemical properties. Hydrogen is the lightest of all known substances (14.4 times lighter than air), and its density is 0.0899 g/liter at 0° C and 1 atmosphere (atm). Hydrogen boils (liquefies) and melts (solidifies) at −252.6° C and −259.1° C, respectively (only helium has lower boiling and melting temperatures). Because the critical temperature is very low (−240° C), liquefaction is difficult. The critical pressure is 12.8 kilograms-force/cm2 (kgf/cm2) or 12.8 atm, and the critical density is 0.0312 g/cm3. Of all the gases hydrogen has the highest heat conductivity, which, at 0° C and 1 atm is 0.174 watt/(m × °K), that is, 4.16 x 10−4 calorie/sec × cm × ° C. The specific heat of hydrogen at 0° C and 1 atm is 14.208 × 103 joules/(kg × ° K), that is, 3.394 calories/(g × ° C). Hydrogen dissolves poorly in water (0.0182 milliliter/g at 20° C and 1 atm) but well in many metals (Ni, Pt, Pd, and other metals), especially in palladium (850 volumes to 1 volume of Pd). Because of its solubility in metals, hydrogen is capable of diffusing through them. The diffusion through a carbon alloy (for example, steel) is sometimes followed by destruction of the alloy as a result of the interaction between hydrogen and carbon (so-called decarbonization). Liquid hydrogen is very light (density at −253° C is 0.0708 g/cm3) and fluid (viscosity at −253° C is 13.8 centipoises).
In most of its compounds hydrogen displays a valence (more accurately, a degree of oxidation) of 1, similar to sodium and other alkali metals; usually hydrogen is considered as an analog of these metals, heading group I of Mendeleev’s system. However, in metallic hydrides the hydrogen ion is charged negatively (the degree of oxidation is −1), that is, the hydride Na+H− is formed similarly to the chloride Na+Cl−. This and some other facts (similar physical properties of H and the halogens, the ability of halogens to replace H in organic compounds) give a basis for attributing hydrogen to group VII of the periodic system as well. Under normal conditions, molecular hydrogen is relatively inactive, combining directly only with the most active metals (fluorine, also chlorine in light). However, when heated, it reacts with many elements. Atomic hydrogen has increased chemical activity in comparison with molecular hydrogen. With oxygen, hydrogen forms water—H2 + ½ O2 = H2O—generating 285.937 × 103 joules/mole—that is, 68.3174 kilocalories/mole of heat (at 25° C and 1 atm). At normal temperatures, this reaction proceeds very slowly, and at temperatures higher than 550° C, with explosions. The explosive limits of a hydrogen-oxygen mixture range from 4 to 94 percent (by volume) of H2 and of a hydrogen-air mixture, from 4 to 74 percent of H2 (a mixture of two volumes of H2 and one volume of O2 is called “detonating gas”). Hydrogen is used in the reduction of many metals because it takes away oxygen from its oxides—for example, CuO + H2 = Cu + H2O and Fe3O4 + 4H2 = 3Fe + 4H2O. With halogens hydrogen forms hydrogen halides—for example, H2 + Cl2 = 2HCl. With fluorine hydrogen explodes (even in the dark and at −252°C; with chlorine and bromium it reacts only in light or when heated and with iodine only when heated. Hydrogen interacts with nitrogen to form ammonia—3H2 + N2 = 2NH3—only when a catalyst is present and at elevated temperature and pressure. When heated, hydrogen reacts vigorously with sulfur—H2 + S = H2S (hydrogen sulfide); it reacts with difficulty with selenium and tellurium. With pure carbon, hydrogen may react without a catalyst only at high temperatures—2H2 + C (amorphous) = CH4 (methane). Hydrogen reacts directly with some metals (alkali, alkaline-earth, and others) forming hydrides—H2 + 2Li = 2LiH. Of important practical value is the reaction between hydrogen and carbon oxide, where, depending on the temperature, pressure and the catalyst, various organic compounds, such as HCHO and CH3OH, are formed. Unsaturated hydrocarbons react with hydrogen, forming saturated hydrocarbons—CnH2n + H2 = CnH2n+2.
Hydrogen and its compounds play a very important role in chemistry. Hydrogen determines the acidic properties of the so-called protonic acids. Hydrogen tends to form the so-called hydrogen bond with some elements. The properties of many organic and inorganic compounds are determined by this bonding.
Preparation. The main raw materials for the production of hydrogen are natural flammable gases, coke gas, and the gases of petroleum refining as well as the products of the gasification of hard and liquid fuel (mainly coal). Hydrogen is also obtained from water by electrolysis (in places where electric energy is cheap). The most important methods of producing hydrogen from natural gases are catalytic interaction of hydrocarbons, mainly methane, with steam (reforming)—CH4 + H2O = CO + 3H2—and a partial oxidation of hydrocarbons with oxygen—CH4 + ½O2 = CO + 2H2. The formed carbon monoxide is also subjected to reforming—CO + H2O = CO2 + H2. The production of hydrogen from natural gases is the least expensive method. A very widespread method of hydrogen production uses water-gas and steam and air gas obtained in coal gasification. The process is based on a conversion of carbon monoxide. Water gas contains up to 50 percent H2 and 40 percent CO. In steam and air gas besides H2 and CO, a significant amount of N2 is found, which is used together with the produced hydrogen for NH3 synthesis. Hydrogen is separated from the coke gas and the refinery gases by removing the other components (they liquefy easier than hydrogen) at very low temperatures. Water electrolysis is conducted by direct current, which is passed through a solution of KOH or NaOH (to avoid corrosion of the steel equipment, acids are not used). In a laboratory, hydrogen is obtained from the electrolysis of water and also from a reaction between zinc and hydrochloric acid. However, the available commercial hydrogen in cylinders is used more often.
Uses. Hydrogen production on the industrial level started at the end of the 18th century, when hydrogen was used for inflating balloons. At the present time hydrogen is widely used in the chemical industry, mainly for the production of ammonia. Hydrogen is also widely used in the production of methyl and other alcohols, synthetic gasoline (synthine), and other products obtained from hydrogen and carbon monoxide by way of synthesis. Hydrogen is also used in the hydrogenation of hard and heavy liquid fuel, fats, and other items, in the synthesis of HCl, in the hydrorefining of petroleum products, in the welding and cutting of metals by oxygen-hydrogen flame (temperatures to 2800° C), and in atomic-hydrogen welding (up to 4000° C). The isotopes deuterium and tritium have important uses in atomic energy.
REFERENCESNekrasov, B. V. Kurs obshchei khimii, 14th ed. Moscow, 1962.
Remy, H. Kurs neorganicheskoi khimii, vol. 1. Moscow, 1963. (Translated from German.)
Egorov, A. P., D. I. Shereshevskii, and I. V. Shmanenkov. Obshchaia khimicheskaia tekhnologiia neorganicheskikh veshchestv, 4th ed. Moscow, 1964.
Obshchaia khimicheskaia tekhnologiia, vol. 1. Edited by S. I. Vol’fkovich. Moscow, 1952.
Lebedev, V. V. Vodorod, ego poluchenie i ispol’zovanie. Moscow, 1958.
Nalbandian, A. B., and V. V. Voevodskii. Mekhanism okisleniia i goreniia vodoroda. Moscow-Leningrad, 1949.
Kratkaia khimicheskaia entsiklopediia, vol. 1. Moscow, 1961. Pages 619-24.
S. E. VAISBERG